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Q: What does H2 Chemistry Notes: Topic 4 - Theories of Acids and Bases cover? A: Compare Arrhenius, Brønsted-Lowry, and Lewis definitions while mastering pH calculations, buffer design, and titration curves for Core Idea 2 (Theories of Acids and Bases).
Acid-base theory appears across Papers 1-4, from conceptual MCQs to practical titrations. This note organises the definitions, equilibrium expressions, and exam-standard calculations required in 2026. Continue exploring linked drills at https://eclatinstitute.sg/blog/h2-chemistry-notes.
Status: SEAB H2 Chemistry (9476, first exam 2026) syllabus and Chemistry Data Booklet last checked 2026-01-13. Core Idea 2 Topic 4 scope is assessed across Papers 1–3.
Quick revision box
What this topic tests: Acid-base models, pH/buffer calculations, and titration curve interpretation.
Top mistakes to avoid: Using wrong acid-base model; buffer equation misuse; poor link between curve features and chemistry.
20-minute sprint plan: 5 min Arrhenius/Brønsted/Lewis contrast; 10 min pH-buffer-titration drill; 5 min curve annotation.
State the theory explicitly in answers. For example, “According to Brønsted-Lowry theory, ammonia acts as a base because it accepts a proton to form NHX4X+.”
Definitions above follow the IUPAC Gold Book. The SEAB Chemistry Data Booklet provides Kw (at 298K) and key constants; Ka/Kb values are provided in the question when needed.
2 Quantifying Acid Strength
2.1 Ka and pKa
For a weak acid HA, the acid dissociation constant is
Ka=[HA][HX+][AX−],pKa=−log10Ka.
Smaller pKa values indicate stronger acids (greater dissociation). Always include state symbols when writing equilibrium equations, for example:
HA(aq)+HX2O(l)HX3OX+(aq)+AX−(aq).
2.2 Polyprotic Acids
Each deprotonation has its own Ka. When calculating pH, consider whether steps are independent. Typically, the first dissociation dominates, and subsequent Ka values are much smaller.
3 pH Calculations
3.1 Strong Acids and Bases
Assume complete dissociation:
Solute
Reaction
pH/pOH shortcut
HCl
HClHX++ClX−
pH=−log10[HX+]
NaOH
NaOHNaX++OHX−
pH+pOH=14 assumes Kw=1.0×10−14 at 298K from the SEAB Chemistry Data Booklet; if the paper supplies a different temperature or Kw, adjust accordingly.
3.2 Weak Acids/Bases
Use equilibrium tables. Example for CH3COOH:
Species
Initial
Change
Equilibrium
CHX3COOH
C
−x
C−x
HX+
0
+x
x
CHX3COOX−
0
Then Ka=C−xx2. If x is much smaller than C, you can approximate C−x≈C to simplify the algebra.
After solving for x, compute pH: pH=−log10x.
Ethanoic acid: weak-acid reference for Ka, ICE tables, and buffer calculations.
The −COOH group is the acidic site, while the stabilised conjugate base CHX3COOX− supports the weak-acid equilibrium model used in Paper 2 calculations.
Use the Ka or Kb values given in the question rather than memorised numbers (they are not listed in the SEAB Chemistry Data Booklet).
4 Buffers and Henderson-Hasselbalch
Buffers resist pH change when small amounts of acid or base are added. For a weak acid buffer:
pH=pKa+log10([HA][AX−])
Use the Henderson-Hasselbalch equation only after confirming both species are present in appreciable amounts. When strong acid/base is added, adjust moles first, then recompute concentrations before applying the equation.
Design principle: Choose pKa close to the desired buffer pH (within ±1) for maximum capacity.
For basic buffers, the same logic applies to the NHX4X+/NHX3 pair: NHX3 uses its nitrogen lone pair to accept HX+, reducing sudden pH swings.
Ammonia: Brønsted-Lowry base and common weak-base buffer component.
5 Titration Curves
Recognise four canonical curves: strong acid vs strong base, strong acid vs weak base, weak acid vs strong base, weak acid vs weak base. Key checkpoints:
Initial pH: Based on identity of the analyte.
Buffering region: Exists when a weak acid/base titrated with strong counterpart.
Half-neutralisation point:pH=pKa (or pOH=pKb).
Equivalence point: Determine species present; e.g. weak acid + strong base yields basic salt solution.
State pH values qualitatively if data not supplied (e.g. “equivalence point pH>7 due to hydrolysis of the conjugate base”).
6 Worked Example
Question:25.0mL of 0.100molL−1 ethanoic acid is titrated with 0.100mol⋅L−1 sodium hydroxide. Ka(CHX3COOH)=1.8×10−5. Calculate the pH after 12.5mL of NaOH has been added.
Solution:
Initial moles: n(CHX3COOH)=CV=(0.100molL−1)(25.0mL)=0.00250mol. Added base moles: n(NaOH)=(0.100molL−1)(12.5mL)=0.00125mol.
Statement: pH=4.74 (rounded to two decimal places).
At half-neutralisation, [AX−]=[HA], so pH=pKa-a result worth memorising.
7 Practical Considerations
Indicators: Choose based on equivalence point pH. For weak acid-strong base titrations, phenolphthalein matches the basic equivalence region.
Use indicator ranges provided in the paper; if a table is supplied, cite the range directly instead of relying on memory.
Paper 4 tasks: When planning buffer preparation, describe stepwise calculations (moles before diluting, then the dilution formula C1V1=C2V2), and specify precise volumetric equipment (pipette, burette).
8 Common Pitfalls
Applying Henderson-Hasselbalch without confirming both conjugate pairs remain.
Forgetting to include the auto-ionisation of water for highly dilute strong acids or bases (e.g. when [HX+]<1×10X−6mol⋅L-1).
Mixing units (molarity vs mole).
Assuming equivalence point pH=7 for all titrations.
9 Drill Suggestions
Rank HClO, HClOX2, HClOX3, and HClOX4 by acid strength; justify using inductive effect and oxidation state.
Design a buffer at pH=9.2 using the NHX4X+/NHX3
Sketch titration curves for (a) HNOX3 vs KOH and (b) CHX3COOH
Common exam mistakes
Assuming equivalence point pH is always 7: This is only true for strong acid–strong base titrations. A weak acid–strong base equivalence point gives pH > 7 due to conjugate base hydrolysis; weak base–strong acid gives pH < 7.
Applying Henderson-Hasselbalch before confirming buffer conditions: The equation is only valid when both the weak acid and its conjugate base are present in appreciable amounts. Using it at or near the equivalence point gives a meaningless result.
Forgetting to adjust moles before using Henderson-Hasselbalch: When strong acid or base is added to a buffer, convert moles and recompute the ratio of conjugate pair first; skipping this step produces a wrong pH.
Mixing up conjugate acid-base pairs: The conjugate base of a weak acid is the species formed after proton donation; identifying the wrong species loses marks in Brønsted-Lowry definition questions.
Not specifying the acid-base model in answers: Writing "acts as a base" without stating "according to Brønsted-Lowry theory" is incomplete; the question usually specifies which model to use.
Ignoring water autoionisation at very low concentrations: For highly dilute strong acid solutions where [HX+]<1⋅10−6mol⋅L−1, water's contribution to [HX+] becomes significant and cannot be ignored.
Choosing an indicator with the wrong pH range: The indicator's transition range must overlap with the equivalence point pH. Methyl orange (range ~3–4.5) is unsuitable for weak acid–strong base titrations.
Frequently asked questions
What is the difference between the Brønsted-Lowry and Lewis definitions of a base? A Brønsted-Lowry base is a proton acceptor, while a Lewis base is an electron-pair donor. All Brønsted-Lowry bases are Lewis bases (lone pair donated to accept H⁺), but Lewis bases also include species that donate electron pairs to non-proton acceptors such as metal ions or BFX3. The Lewis definition is broader and covers complex formation and coordination chemistry.
How do I know which acid-base definition to use in a question? Apply the definition that the question specifies or the one that fits the chemistry described. If water is the solvent and proton transfer is occurring, Brønsted-Lowry is appropriate. If the question involves electron pair donation to a metal ion or a Lewis acid like BFX3, use the Lewis definition.
What does pH = pKa at the half-neutralisation point mean in practice? At half-neutralisation, exactly half the weak acid has been converted to its conjugate base, so [HA]=[AX−] and the log term in Henderson-Hasselbalch equals zero. This is a useful shortcut for reading pKa directly from a titration curve at the midpoint volume.
Why does the pH change slowly in the buffering region of a titration curve? In the buffering region, both the weak acid and its conjugate base are present in significant amounts. Adding small amounts of strong base converts some HA to AX− but the ratio changes only slightly, so pH changes very little - this is the definition of buffer action.
Struggling with Theories of Acids and Bases? Our H2 Chemistry tuition programme covers this topic with structured practice, Paper 4 practical drills, and worked exam solutions.
